Electronic Configuration Of Elements

Edited By Shivani Poonia | Updated on Jul 02, 2025 05:53 PM IST

Introduction to Electronic Configuration of Elements in Chemistry

Knowledge of electronic configuration in elements is central to chemistry since it acts as a roadmap for properties/behaviour and reactivity of elements. Atoms are made up of a nucleus with protons neutrons and electrons that revolve around the nucleus in orbits known as shells. These electrons are placed in a manner that obeys certain regulations as constrained by quantum mechanics known as the Pauli exclusion principle, Hund’s rule and the Aufbau rule. The electronic configuration of an element defines how the respective element possesses electrons in the mentioned energy levels or orbitals defined by the quantum numbers; the principal quantum number (n), the azimuthal quantum number (l), the magnetic quantum number(ml), and the spin quantum number (ms).This allocation not only defines its chemical activity and connection abilities but also describes its placement and functioning scope within a periodical table.

This Story also Contains

Electronic Configuration of Elements-
Solved Examples Based On Electronic Configuration of Elements
Conclusion

Electronic Configuration Of Elements

In this article, we will cover the concept of Electronic Configuration of Elements. This concept falls under the broader category of Atomic structure, which is a crucial chapter in Class 11 chemistry. It is not only essential for board exams but also for competitive exams like the Joint Entrance Examination (JEE Main), National Eligibility Entrance Test (NEET), and other entrance exams such as SRMJEE, BITSAT, WBJEE, BCECE and more.

Let us study in detail the Electronic Configuration of Elements to gain insights into this topic and solve a few related problems.

Electronic Configuration of Elements-

Writing of electronic configuration of any element is based on three rules as told in previous concepts. They are:

Aufbau principle
Pauli’s exclusion principle
Hund’s rule of maximum multiplicity

The distribution of electrons into orbitals of an atom is called its electronic configuration.

The electronic configuration of the different elements can be represented in two ways:

Subshell notation
Orbital diagram

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Let’s understand with the help of examples:

The hydrogen atom has only one electron which goes into the orbital with the lowest energy, namely 1s. The E.C of hydrogen is 1s¹ (subshell notation)

Orbital diagram:

2. The electronic configuration of lithium is 1s² 2s¹. Lithium has 3 electrons, as its atomic number is 3. 2 electrons filled in 1s orbital and 1 electron filled in 2s orbital.

let us consider fluorine (Z = 9) :

F(Z = 9) = 1s², 2s², 2px², 2py², 2pz¹ or

The importance of knowing the exact electronic configuration of an element lies in the fact that the chemical properties of an element are dependent on the behaviour and relative arrangement of its electrons.

Solved Examples Based On Electronic Configuration of Elements

Example 1: The electronic configuration of copper is:

1) (correct) [Ar]3d¹⁰4s¹

2) [Ar]3d⁹4s²

3) [Ar]3d¹⁰4s²

4) [Ar]3d⁸4s²

Solution: Ideally, the electronic configuration of Cu must be [Ar] 3d⁹ 4s² but in this case, the electrons in d-orbitals are not symmetrically filled. Thus to maintain the symmetricity, one electron from the 4s-orbital goes to the d-orbital and thus Cu maintains the electronic configuration as [Ar] 3d¹⁰ 4s¹.

Hence, the answer is the option (1)

Example 2: Which law indicates the pairing of electrons in the same orbital?

1) Newton’s first law

2) (correct) Hund’s rule

3) Aufbau principle

4) Pauli exclusion principle

Solution: Hund’s rule states that “pairing of electrons in the orbitals belonging to the same subshell (p, d or f) does not take place until each orbital belonging to that subshell has got one electron each. It is singly occupied”.

Hence, the answer is the option (2).

Example 3: In which of the following pairs, the outermost electronic configuration will be the same?

1) (correct) Cr⁺and Mn²⁺
2) V²⁺ and Cr⁺
3) Ni²⁺ and Cu⁺
4) Fe²⁺ and Co⁺

Solution: The outermost electronic configurations-
Cr+→[Ar]3d⁵Mn²⁺→[Ar]3d⁵
V2+→[Ar]3d³Cr⁺→[Ar]3d⁵
Ni2+→[Ar]3d⁸Cu⁺→[Ar]3d¹⁰

In Cr+and Mn2+, the outermost electronic configuration is the same.

Hence, the answer is the option (2).

Example 3: In the sixth period, the orbitals that are filled are :

1) (correct) 6s, 4f, 5d, 6p

2) 6s, 5d, 5f, 6p

3) 6s, 5f, 6d, 6p

4) 6s, 6p, 6d, 6f

Solution: Energy order of orbitals according to Aufbau principle-

The order of orbitals filling is 6s, 4f, 5d, 6p.

Therefore, the correct option is (1).

Example 4: Element "E" belongs to the period 4 and group 16 of the periodic table. The valence shell electron configuration of the element, which is just above "E" in the group is

1) (correct) 3s²,3p⁴
2) 3d¹⁰,4s²,4p⁴
3) 4d¹⁰,5s²,5p⁴
4) 2s²,2p⁴

Solution:
The element " E " will be Se (Selenium), it belongs to period 4 and group 16 of the periodic table.
Just above Se, S (sulfur) is present. The valence shell electronic configuration of Sulfur is 3s²,3p⁴
Hence, the answer is the option (1).

Example 5: The electronic configuration of Pt (atomic number 78) is :
1) (correct) [Xe]4f¹⁴5 d⁹6s¹
2) [Kr]4f¹⁴5 d¹⁰
3) [Xe]4f¹⁴5 d¹⁰
4) [Xe]4f¹⁴5 d⁸6 s²

Solution:

The electronic configuration of Pt(78) is [Xe]4f¹⁴5 d⁹6 s¹

Hence, the answer is the option (1).

Conclusion

All in all, the study of electronic configurations deepens one’s understanding of how atoms are built and how this shapes the chemical world. On this basis, we eventually explain the reactivity, bonding, as well as behaviours of different elements in the periodic table by understanding the distribution of electrons’ energy levels in the atomic structures. Pauli exclusion principle states that no two electrons present in an atom can have the same four quantum numbers and Hund’s rule for establishing the electron configuration to minimize energy is critical in the chemists’ attempts to predict the chemical properties of an element and its reactivity with other substances.

Frequently Asked Questions (FAQs)

1. How do you define electronic configuration?

Electronic configuration in other words is referred to as the arrangement of electrons in an atom about the nucleus based on quantum numbers.

2. How does one determine the e configuration of an element?

It follows the Aufbau principle which involves the filling of orbitals followed by electrons Pauli’s exclusion principle for each orbital; and Hund’s rule for degenerate orbitals.

3. What is the meaning of Electronic Configuration?

It expresses where an element is situated on the table of the periodic and other chemical characteristics like reactivity and in bonding.

4. What are valence electrons and how are they determined from electronic configuration?

Valence electrons are the electrons in the outermost shell of an atom, which participate in chemical bonding. They can be determined by looking at the electrons in the highest energy level (largest n value) in the electronic configuration.

5. What is the Aufbau principle?

The Aufbau principle states that electrons fill orbitals from lowest energy to highest energy. This means that lower energy levels are filled before higher ones, following the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on.

6. How does electronic configuration explain the formation of ions?

Electronic configuration explains ion formation by showing how atoms gain or lose electrons to achieve a stable noble gas configuration. For example, sodium (2,8,1) loses one electron to become Na+ (2,8), while chlorine (2,8,7) gains one electron to become Cl- (2,8,8).

7. What is the aufbau principle, and are there any exceptions to it?

The aufbau principle states that electrons fill orbitals from lowest to highest energy. However, there are exceptions, particularly in transition metals, where factors like electron-electron repulsion can cause electrons to fill higher energy orbitals before lower ones are completely filled.

8. What is the significance of the lanthanide and actinide series in terms of electronic configuration?

The lanthanide and actinide series involve the filling of f-orbitals, which are buried deep within the electron cloud. This leads to similar chemical properties across each series and unique phenomena like the lanthanide contraction.

9. What is electronic configuration?

Electronic configuration is the arrangement of electrons in an atom's orbitals. It describes how electrons are distributed in different energy levels and sublevels around the nucleus.

10. Why is understanding electronic configuration important?

Understanding electronic configuration is crucial because it determines an element's chemical properties, reactivity, and bonding behavior. It helps explain periodic trends and predict how elements will interact with each other.

11. How does the Pauli exclusion principle relate to electronic configuration?

The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers. This means that each orbital can hold a maximum of two electrons with opposite spins, influencing how electrons are arranged in an atom.

12. What is Hund's rule and how does it apply to electronic configuration?

Hund's rule states that electrons in orbitals of the same energy (such as p or d orbitals) will occupy separate orbitals before pairing up. This maximizes the total spin and minimizes electron repulsion, leading to a more stable configuration.

13. How do you write the electronic configuration for hydrogen?

The electronic configuration for hydrogen is 1s¹. This means it has one electron in the 1s orbital, which is the lowest energy level.

14. What is the difference between ground state and excited state electronic configurations?

The ground state is the lowest energy configuration of an atom, where electrons occupy the lowest available energy levels. An excited state occurs when an electron is promoted to a higher energy level, resulting in a less stable but more reactive configuration.

15. How does the octet rule relate to electronic configuration?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons (like noble gases). This rule is based on the stability of filled s and p orbitals in the outermost energy level.

16. What is the significance of noble gas electronic configurations?

Noble gas electronic configurations (e.g., 2,8 or 2,8,8) are particularly stable because they have completely filled outer shells. Other elements often gain, lose, or share electrons to achieve similar configurations, which drives many chemical reactions.

17. How does the periodic table reflect electronic configuration?

The periodic table is organized based on electronic configuration. Elements in the same group (column) have similar outer shell configurations, while periods (rows) represent the filling of new energy levels. This organization explains periodic trends in properties.

18. What is the difference between core electrons and valence electrons?

Core electrons are those in completely filled inner shells, while valence electrons are in the outermost, partially filled shell. Core electrons are generally not involved in chemical bonding, whereas valence electrons determine an element's chemical properties.

19. How do you determine the number of unpaired electrons from an electronic configuration?

To determine the number of unpaired electrons, write out the electron configuration and apply Hund's rule. Count the number of single electrons in the highest energy orbitals. For example, nitrogen (1s² 2s² 2p³) has three unpaired electrons in its 2p orbitals.

20. What is electron shielding and how does it affect electronic configuration?

Electron shielding occurs when inner electrons partially block the nuclear charge felt by outer electrons. This effect causes outer electrons to be held less tightly, influencing atomic size, ionization energy, and other properties across the periodic table.

21. How does electronic configuration explain the formation of chemical bonds?

Electronic configuration explains bond formation by showing how atoms share, gain, or lose electrons to achieve stable configurations. For example, covalent bonds form when atoms share electrons to complete their outer shells, while ionic bonds result from electron transfer.

22. What is the relationship between electronic configuration and an element's position in the periodic table?

An element's position in the periodic table directly relates to its electronic configuration. The group (column) number typically corresponds to the number of valence electrons, while the period (row) number indicates the highest occupied energy level.

23. How do you write electronic configurations using noble gas notation?

Noble gas notation simplifies writing electronic configurations by using the symbol of the previous noble gas in brackets, followed by the remaining electron configuration. For example, sodium can be written as [Ne]3s¹ instead of 1s²2s²2p⁶3s¹.

24. What are subshells and how do they relate to electronic configuration?

Subshells (s, p, d, f) are subdivisions of main energy levels that have different shapes and energies. In electronic configuration, they determine the order in which orbitals are filled and influence an element's properties and behavior in chemical reactions.

25. How does the concept of effective nuclear charge relate to electronic configuration?

Effective nuclear charge is the net positive charge experienced by an electron, considering the shielding effect of other electrons. It increases across a period as more protons are added, affecting properties like atomic radius and ionization energy.

26. What is the significance of half-filled and fully-filled subshells in electronic configuration?

Half-filled (e.g., p³, d⁵) and fully-filled (e.g., p⁶, d¹⁰) subshells are often more stable due to symmetrical electron distribution and reduced electron repulsion. This can lead to unexpected electronic configurations in some transition elements.

27. How does electronic configuration explain the formation of colored compounds in transition metals?

The partially filled d-orbitals in transition metal electronic configurations allow for electron transitions between energy levels when exposed to light. These transitions absorb specific wavelengths, resulting in the complementary color being observed.

28. How does spin pairing energy affect electronic configuration?

Spin pairing energy is the energy required to place two electrons with opposite spins in the same orbital. In some cases, particularly in transition metals, it can be more energetically favorable to promote an electron to a higher energy orbital rather than pair it in a lower one.

29. What is the relationship between electronic configuration and magnetic properties of elements?

The number of unpaired electrons in an atom's electronic configuration determines its magnetic properties. Elements with unpaired electrons are paramagnetic and attracted to magnetic fields, while those with all paired electrons are diamagnetic and weakly repelled.

30. How does electronic configuration explain the periodic trends in atomic radius?

Atomic radius generally decreases across a period due to increasing nuclear charge and electron shielding remaining relatively constant. Down a group, atomic radius increases as new energy levels are added, outweighing the effect of increased nuclear charge.

31. What is the connection between electronic configuration and ionization energy?

Ionization energy generally increases across a period due to increasing nuclear charge and decreases down a group due to increased shielding and distance from the nucleus. Fully filled and half-filled subshells also contribute to higher ionization energies.

32. How does electronic configuration relate to an element's electronegativity?

Electronegativity generally increases across a period and decreases down a group, following trends in effective nuclear charge. Elements with nearly filled outer shells (like halogens) tend to have high electronegativity due to their strong attraction for electrons.

33. What is the significance of the 18-electron rule in transition metal complexes?

The 18-electron rule states that stable transition metal complexes often have a total of 18 valence electrons (including those from ligands). This rule is based on filling all available valence orbitals (one s, three p, and five d) for maximum stability.

34. How does electronic configuration explain the formation of molecular orbitals?

Electronic configuration of atoms provides the basis for molecular orbital theory. When atoms bond, their atomic orbitals combine to form molecular orbitals, which can be bonding (lower energy) or antibonding (higher energy), determining the stability of the molecule.

35. What is the relationship between electronic configuration and chemical reactivity?

Chemical reactivity is largely determined by valence electrons. Elements with nearly empty or nearly full valence shells tend to be more reactive as they can easily lose or gain electrons to achieve stable configurations. Noble gases, with full outer shells, are least reactive.

36. How does electronic configuration explain the formation of coordinate covalent bonds?

In coordinate covalent bonds, one atom (the donor) provides both electrons for the bond. This is possible when one atom has a lone pair of electrons in its valence shell that can be shared with another atom (the acceptor) that has an empty orbital to accommodate them.

37. What is the significance of the lanthanide contraction in electronic configuration?

The lanthanide contraction refers to the smaller-than-expected increase in atomic radius across the lanthanide series. This is due to poor shielding by 4f electrons, leading to increased effective nuclear charge and a contraction of the electron cloud.

38. How does electronic configuration explain the difference between diamagnetic and paramagnetic substances?

Diamagnetic substances have all paired electrons in their electronic configuration, resulting in no net magnetic moment. Paramagnetic substances have one or more unpaired electrons, creating a net magnetic moment that causes attraction to magnetic fields.

39. What is the relationship between electronic configuration and spectroscopic properties of elements?

Electronic configuration determines the possible electron transitions within an atom. These transitions correspond to specific energy differences, which are observed as distinct spectral lines when atoms absorb or emit light, forming the basis of spectroscopy.

40. How does the concept of hybridization relate to electronic configuration?

Hybridization involves the mixing of atomic orbitals to form new hybrid orbitals of equal energy. This process is based on the electronic configuration of an atom and explains molecular geometries that cannot be accounted for by simple atomic orbital theory.

41. What is the significance of the inert pair effect in electronic configuration?

The inert pair effect refers to the tendency of the outermost s electrons in some heavy elements to resist participation in bonding. This is due to increased stability of the s orbital and poor shielding of the nuclear charge by d and f electrons.

42. How does electronic configuration explain the formation of complex ions?

Complex ions form when a central metal atom or ion bonds with surrounding ligands. The electronic configuration of the metal determines the number and type of ligands it can accommodate, as well as the geometry of the resulting complex.

43. What is the relationship between electronic configuration and photoelectric effect?

The photoelectric effect occurs when light causes electrons to be ejected from a material. The electronic configuration determines the binding energy of electrons, which in turn determines the minimum frequency of light required to cause electron ejection.

44. How does electronic configuration relate to the concept of electron affinity?

Electron affinity is the energy change when an atom in the gas phase gains an electron. It is generally more exothermic for elements with nearly full valence shells, as they can achieve stable noble gas configurations by gaining electrons.

45. What is the significance of Slater's rules in understanding electronic configuration?

Slater's rules provide a method for estimating the effective nuclear charge experienced by electrons in multi-electron atoms. These rules consider the shielding effects of other electrons and help explain trends in atomic properties across the periodic table.

46. How does electronic configuration explain the formation of metallic bonds?

Metallic bonds form when metal atoms share their valence electrons in a "sea of electrons." This is possible because metals typically have few valence electrons and low ionization energies, allowing these electrons to be delocalized across the entire structure.

47. What is the relationship between electronic configuration and crystal field theory?

Crystal field theory explains the behavior of transition metal complexes based on the splitting of d-orbitals in different ligand environments. The electronic configuration of the metal ion determines how these d-orbitals are filled and thus affects the complex's properties.

48. How does electronic configuration relate to the concept of electronegativity?

Electronegativity is influenced by the effective nuclear charge and the distance of valence electrons from the nucleus. Elements with nearly full valence shells and small atomic radii tend to have high electronegativity due to their strong attraction for electrons.

49. What is the significance of the aufbau principle in predicting chemical properties?

The aufbau principle helps predict chemical properties by determining the order in which orbitals are filled. This information is crucial for understanding valence electrons, which largely determine an element's chemical behavior and reactivity.

50. How does electronic configuration explain the formation of sigma and pi bonds?

Sigma bonds form through head-on overlap of atomic orbitals, while pi bonds form through side-by-side overlap. The type of bond formed depends on the electronic configuration and the available orbitals of the atoms involved in bonding.

51. What is the relationship between electronic configuration and the periodic trends in electron affinity?

Electron affinity generally increases across a period due to increasing nuclear charge and decreases down a group due to increasing atomic size. Exceptions occur for noble gases (stable configurations) and some elements with half-filled subshells.

52. How does electronic configuration relate to the concept of ionization energy?

Ionization energy is the energy required to remove an electron from an atom. It is influenced by factors such as nuclear charge, electron shielding, and the stability of certain electronic configurations (e.g., noble gas configurations).

53. How does electronic configuration explain the formation of hydrogen bonds?

Hydrogen bonds form between a hydrogen atom bonded to a highly electronegative atom (like N, O, or F) and another highly electronegative atom. The electronic configuration of these elements results in a partial positive charge on the hydrogen, enabling this type of bonding.